Question

Following are several Lewis structures showing all valence electrons. Assign formal charges in each structure as appropriate. A. B. C. D. E. F.

Short answer

Question: Calculate the formal charges for each atom in the Lewis structure of nitrate (NO3-) ion and assign them to the corresponding atoms. Step 1: Understanding Formal Charge Recall the formula for calculating formal charge: (Formal Charge) = (Valence electrons of the atom) - (Non-bonding electrons) - 1/2(Bonding electrons) Step 2: Identifying Valence Electrons, Non-Bonding Electrons, and Bonding Electrons In the nitrate ion (NO3-), we have the following valence electrons for each atom: - Nitrogen (N): 5 valence electrons. - Oxygen (O): 6 valence electrons. Since there are three oxygen atoms in the nitrate ion, we will calculate the formal charge for one oxygen atom and multiply it by three. Step 3: Calculating Formal Charge for Each Atom For Nitrogen (N) atom in NO3-: Valence electrons = 5 Non-bonding electrons = 0 Bonding electrons = 4 (2 from each single bond, total 2 single bonds) Formal Charge (N) = 5 - 0 - 1/2 (4) = 5 - 0 - 2 = +1 For Oxygen (O) atom in NO3-: Valence electrons = 6 Non-bonding electrons = 6 (3 lone pairs) Bonding electrons = 2 (1 single bond) Formal Charge (O) = 6 - 6 - 1/2 (2) = 6 - 6 - 1 = -1 Since there are three oxygen atoms, the total formal charge for all oxygens is -1 x 3 = -3. Step 4: Assigning Formal Charges in the Lewis Structure In the Lewis structure of the nitrate ion (NO3-), the formal charges are: - Nitrogen (N): +1 - Oxygen (O): -1 for each oxygen atom The overall charge of the nitrate ion is -1, which is the sum of the formal charges of all its constituent atoms: +1 (N) + (-1 x 3 O) = -1.

Step-by-step solution

Understanding Formal Charge

Formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. It is used to find the most likely Lewis structure for a molecule, which will be the one with the least number of formal charges. The formal charge can be calculated using the formula: (Formal Charge) = (Valence electrons of the atom) - (Non-bonding electrons) - 1/2(Bonding electrons)

Identifying Valence Electrons, Non-Bonding Electrons, and Bonding Electrons

To calculate the formal charge of an atom, first, identify the valence electrons, non-bonding electrons (lone pairs), and bonding electrons in the atom. -Valence electrons: The number of electrons in the outermost shell of an atom as indicated in the periodic table. -Non-bonding electrons: The number of electrons in lone pairs on the atom. -Bonding electrons: The number of electrons shared between two atoms in a chemical bond (represented by a single, double, or triple bond).

Calculating Formal Charge for Each Atom

Use the formula mentioned in Step 1 to calculate the formal charge for each atom in the given Lewis structure. Remember to count the bonding electrons correctly depending on the type of bond (single, double, or triple).

Assigning Formal Charges in the Lewis Structure

After calculating the formal charges, assign them to the corresponding atoms in the Lewis structure. Apply these steps to each Lewis structure given in the exercise to assign the appropriate formal charges to each atom.

Key concepts

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom. These are the particles that engage in chemical bonding and are identified by referring to the group number in the periodic table for main group elements. For instance, group 1 elements have 1 valence electron while group 17 elements have 7 valence electrons. Understanding the number of valence electrons an atom has is crucial because it determines how an atom will bond with others, and it is the starting point for drawing a Lewis structure.

When you begin analyzing a Lewis structure, the first thing you do is count the valence electrons for each atom based on its group number. This will guide you in predicting the number and types of bonds an atom can form to achieve a full octet, or eight electrons, reflecting a stable configuration similar to the noble gases.

Non-Bonding Electrons

Non-bonding electrons, also referred to as lone pairs, are valence electrons that are not involved in forming bonds. These electrons are depicted as pairs of dots in the Lewis structure and are localized on one atom. It's important to locate these electrons because they contribute to the formal charge calculation and can affect the shape and polarity of the molecule.

In a Lewis structure, non-bonding electrons are represented by pairs of dots around an atom. For example, the oxygen atom in water (H₂O) has two pairs of non-bonding electrons. It's important to note that non-bonding electrons can influence the physical properties of molecules, such as their boiling point and solubility, because they can create regions of negative electron density that affect how molecules interact with each other.

Bonding Electrons

Bonding electrons are the valence electrons that are shared between atoms to form chemical bonds. They are represented in Lewis structures by lines where a single line indicates a single bond with two electrons, a double line for a double bond with four electrons, and a triple line for a triple bond with six electrons.

For instance, in a molecule of nitrogen gas (N₂), the two nitrogen atoms share six electrons to form a triple bond, represented by three lines between them. In the calculation of formal charge, bonding electrons are counted differently: Rather than counting the full number of bonding electrons around an atom, you count one electron for each bond because the electrons are shared between two atoms.

Lewis Structure Analysis

Lewis structure analysis involves drawing all the possible connections between atoms within a molecule to predict its shape and the distribution of valence electrons. This process includes calculating the formal charges to determine the most stable structure, which typically has formal charges close to zero.

When you analyze a Lewis structure, first count the total number of valence electrons available from all atoms. Then, use lines and dots to denote bonds and lone pairs, and make sure the total count of electrons matches what you started with. After drawing potential structures, calculate the formal charges to assess which structure is most likely. A correct Lewis structure will have minimal formal charges and follow the octet rule when possible, although there are exceptions like in the case of molecules with an odd number of electrons, or those that have atoms capable of having more or fewer than eight electrons.

Chemical Bonds

Chemical bonds are the attractive forces that hold atoms together in molecules. The main types of chemical bonds include ionic, covalent, and metallic bonds. In the context of Lewis structures, we focus on covalent bonds, where atoms share valence electrons to achieve stability.

The formation of a chemical bond results in a molecule with shared bonding electrons, and potentially non-bonding electrons, which can result in a more complicated electron distribution and formal charges on individual atoms. Understanding the types of chemical bonds and how they form is critical for predicting molecular structure, reactivity, and properties. By analyzing Lewis structures and assigning formal charges, we can get a better idea of how molecules will behave in chemical reactions based on their electronic structures.